Welcome to Atomic Structure and the Periodic Table!
Hello future chemists! This chapter is the absolute foundation of everything we will study in bonding and reactions. Think of it as learning the alphabet before you can write a novel. Understanding how atoms are built and how the periodic table is structured gives us the key to predicting chemical behaviour.
Don’t worry if some ideas seem abstract—we will break them down using clear steps and relatable examples. By the end of these notes, you’ll be a pro at understanding the structure that governs all matter!
1. The Building Blocks of the Atom
1.1. Subatomic Particles
Atoms are made of three fundamental particles. It’s vital to know their location, relative mass, and charge.
| Particle | Location | Relative Mass (approx.) | Relative Charge |
|---|---|---|---|
| Proton | Nucleus | 1 | +1 |
| Neutron | Nucleus | 1 | 0 (Neutral) |
| Electron | Orbitals/Shells | 1/1840 (or negligible) | -1 |
Quick Fact: In any neutral atom, the number of protons must equal the number of electrons, so the charges cancel out!
1.2. Defining the Atom: Z and A
We use two key numbers to identify an element:
- Atomic Number (Z): This is the number of protons in the nucleus. (Z is the element's identity card. If Z changes, the element changes!)
- Mass Number (A): This is the total number of particles in the nucleus (protons + neutrons).
Finding Neutrons:
Number of Neutrons = Mass Number (A) – Atomic Number (Z)
Example: Carbon has Z=6 and A=12. It has 6 protons, 6 electrons, and \(12 - 6 = 6\) neutrons.
1.3. Isotopes
Isotopes are atoms of the same element (meaning they have the same number of protons, Z) but have different numbers of neutrons (meaning they have different mass numbers, A).
Analogy: Think of isotopes as different models of the same car—same engine (protons/electrons), but different weights (due to extra neutrons).
Key Takeaway for Section 1: Protons define the element (Z). Neutrons affect the mass (A). Electrons balance the charge.
2. Measuring Mass and Abundance
2.1. Relative Atomic Mass (\(A_r\))
Because most elements exist naturally as a mixture of isotopes, we must calculate an average mass. This average is called the Relative Atomic Mass (\(A_r\)).
The calculation is a weighted average based on the natural abundance of each isotope.
Step-by-Step Calculation:
\[ A_r = \frac{\sum (\text{Isotope Mass} \times \text{Isotope Abundance})}{100} \]
(If abundance is given as a percentage.)
Example: Chlorine exists as two isotopes: Chlorine-35 (75% abundance) and Chlorine-37 (25% abundance).
\[ A_r = \frac{(35 \times 75) + (37 \times 25)}{100} = 35.5 \]
This is why you often see non-whole number masses on the Periodic Table!
2.2. Mass Spectrometry (MS)
Mass spectrometry is the experimental technique used to determine the mass and abundance of isotopes.
Don't worry if this process seems complicated—you need to know the *stages* and what the resulting graph *means*.
The 5 Key Stages of Mass Spectrometry:
- Vaporisation: The sample is turned into a gas.
- Ionisation: The gas atoms are bombarded with high-energy electrons, knocking electrons off the atoms to form positive ions (usually +1 charge). (Only charged particles can be manipulated by electric and magnetic fields.)
- Acceleration: The positive ions are accelerated by an electric field so they all have the same kinetic energy.
- Deflection: The ions pass through a magnetic field. Lighter ions and ions with higher charge are deflected *more* strongly than heavier ions. Analogy: A strong wind (magnetic field) deflects a table tennis ball (light ion) more easily than a bowling ball (heavy ion).
- Detection: The ions hit a detector, which measures the arrival time/intensity. This gives us the mass/charge ratio (m/z) and the relative abundance.
Interpreting the Mass Spectrum
The spectrum shows a series of peaks:
- The position of the peak on the x-axis gives the \(m/z\) ratio (which usually corresponds directly to the isotope's mass, assuming a +1 charge).
- The height of the peak gives the relative abundance of that isotope.
Common Mistake to Avoid: When calculating \(A_r\) from a spectrum, make sure your percentages (relative heights) add up to 100% before using the weighted average formula!
Key Takeaway for Section 2: Mass spectrometry sorts atoms by mass to determine the exact percentages needed to calculate the weighted average relative atomic mass (\(A_r\)).
3. Electron Structure and Configuration
3.1. Shells and Subshells
Electrons are not randomly scattered; they exist in specific energy levels (shells). Within each shell are specific regions of space called subshells (or sublevels).
- Shells (Principal Quantum Numbers, n=1, 2, 3...): Higher the number, higher the energy.
- Subshells: Within the shells are different types of subshells: s, p, d, and f.
| Shell (n) | Subshells Present | Total Orbitals | Max Electrons |
|---|---|---|---|
| 1 | s | 1 | 2 |
| 2 | s, p | 1 + 3 = 4 | 8 |
| 3 | s, p, d | 1 + 3 + 5 = 9 | 18 |
3.2. Orbitals and Shapes
A subshell is made up of one or more orbitals. An orbital is a specific region of space that can hold a maximum of two electrons.
- s-orbital: Spherical shape. There is 1 s-orbital per subshell.
- p-orbital: Dumbbell shape. There are 3 p-orbitals (\(p_x\), \(p_y\), \(p_z\)) per subshell, aligned along the axes.
- d-orbital: More complex shapes (5 d-orbitals per subshell).
3.3. Rules for Filling Electrons
Electrons fill orbitals according to three main rules:
-
The Aufbau Principle (Building Up): Electrons fill the lowest energy levels first.
The filling order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...
(Don’t panic about 4s filling before 3d. This is just the way the energy levels overlap!)
- Pauli Exclusion Principle: An orbital can hold a maximum of two electrons, and these must have opposite spin (represented by arrows pointing up and down).
- Hund’s Rule (Bus Seat Rule): When filling degenerate orbitals (orbitals with the same energy, like the three p-orbitals), electrons will fill them singly first before pairing up. Analogy: When you get on a bus, you try to sit alone before you sit next to someone!
Example: Oxygen (Z=8)
1. Start filling: 1s goes first (holds 2e). -> \(1s^2\)
2. Next is 2s (holds 2e). -> \(2s^2\)
3. Remaining 4 electrons go into 2p. (2p can hold 6e). -> \(2p^4\)
Electron Configuration: \(1s^2 2s^2 2p^4\)
Quick Review: The superscript number tells you how many electrons are in that subshell.
Key Takeaway for Section 3: Electrons fill defined orbitals (s, p, d) following specific rules (low energy first, two per orbital, fill singly before pairing) to achieve the most stable arrangement.
4. Ionization Energy: Peeling the Electron Layers
4.1. Definition of First Ionization Energy (\(IE_1\))
Ionization Energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
It must be done in the gaseous state because we want to measure the energy required to break the nucleus-electron attraction, not the energy required to break lattice forces (in solids) or intermolecular forces (in liquids).
General Equation:
\[ X(g) \rightarrow X^+(g) + e^- \quad (\Delta H = IE_1) \]
Ionization energy is always endothermic (positive \(\Delta H\)) because energy must be supplied to overcome the attraction between the nucleus and the electron.
4.2. Factors Affecting Ionization Energy
The easier it is to remove an electron, the lower the IE. This depends on three factors:
- Nuclear Charge (Number of Protons): Higher positive charge pulls electrons in more tightly, increasing IE.
- Atomic Radius/Distance: If the valence electron is further from the nucleus, the attraction is weaker, decreasing IE.
- Shielding: Inner shell electrons "shield" the outer electrons from the full attractive force of the nucleus. More shielding means lower IE.
Mnemonic: Remember the three factors as CND (Charge, iNcreasing Distance, Shielding).
4.3. Successive Ionization Energies
You can remove more than one electron, leading to the second, third, and successive ionization energies (\(IE_2, IE_3\), etc.).
\[ X^+(g) \rightarrow X^{2+}(g) + e^- \quad (\Delta H = IE_2) \]
Successive IEs always increase because you are removing an electron from an increasingly positive ion, meaning the remaining electrons are held tighter.
Evidence for Shell Structure
Plotting successive IEs reveals massive jumps in energy. These jumps occur when an electron is removed from a shell closer to the nucleus (i.e., a shell with lower shielding and smaller radius).
Analogy: Imagine peeling layers from an onion. The outer layers come off easily (low IE). Once you hit the core (the next principal quantum shell), it’s much harder to peel (massive jump in IE).
The number of electrons in each shell can be determined by counting the number of IEs between the large jumps.
Key Takeaway for Section 4: Ionization energy measures how hard it is to remove an electron. Huge jumps in successive IEs prove that electrons are arranged in distinct shells.
5. The Periodic Table: Organisation and Structure
The modern periodic table is arranged by increasing atomic number (Z).
5.1. Groups and Periods
- Periods (Horizontal Rows): Indicate the principal quantum shell (energy level) that the outermost electrons occupy. (E.g., Elements in Period 3 start filling the n=3 shell).
- Groups (Vertical Columns): Elements in the same group have the same number of outer shell (valence) electrons, leading to similar chemical properties. (E.g., Group 2 elements all have \(ns^2\) configuration).
5.2. Blocks of the Periodic Table
The table is divided into blocks corresponding to the type of subshell being filled:
- s-block: Groups 1 and 2 (plus Helium).
- p-block: Groups 13 to 18.
- d-block: The Transition Metals (Groups 3 to 12).
- f-block: The Lanthanides and Actinides.
5.3. Brief Overview of Periodic Trends
Understanding electron configuration allows us to explain the change in physical and chemical properties across periods and down groups.
Trend Across a Period (e.g., Period 3: Na to Ar)
- Nuclear Charge: Increases (more protons).
- Shielding: Stays roughly constant (electrons are all in the same principal shell).
- Atomic Radius: Decreases. Increased nuclear attraction pulls the constant number of electron shells closer.
- First Ionization Energy: Generally increases. Higher nuclear charge pulls electrons tighter, making them harder to remove. (There are small dips at the start of the p-block and when Hund’s rule requires electron pairing, but the overall trend is up).
Trend Down a Group (e.g., Group 2: Be to Ba)
- Number of Outer Electrons: Stays constant (e.g., all have 2).
- Atomic Radius: Increases. New principal quantum shells are added, increasing the distance from the nucleus.
- Shielding: Increases significantly (more inner shells).
- First Ionization Energy: Decreases. Increased distance and shielding overcome the increased nuclear charge, making the outer electron easier to remove.
Did you know? The existence and properties of elements like Gallium (Ga) and Germanium (Ge) were predicted by Mendeleev decades before they were discovered, simply based on the gaps in his periodic table structure!
Key Takeaway for Section 5: Group number tells you valence electrons; Period number tells you the highest occupied shell. Trends across the table are governed by the balance between nuclear charge, shielding, and distance.